======================================================================
= Iodine =
======================================================================
Introduction
======================================================================
Iodine is a chemical element; it has symbol I and atomic number 53.
The heaviest of the stable halogens, it exists at standard conditions
as a semi-lustrous, non-metallic solid that melts to form a deep
violet liquid at 114 C, and boils to a violet gas at 184 C. The
element was discovered by the French chemist Bernard Courtois in 1811
and was named two years later by Joseph Louis Gay-Lussac, after the
Ancient Greek , meaning 'violet'.
Iodine occurs in many oxidation states, including iodide (I−), iodate
(), and the various periodate anions. As the heaviest essential
mineral nutrient, iodine is required for the synthesis of thyroid
hormones. Iodine deficiency affects about two billion people and is
the leading preventable cause of intellectual disabilities.
The dominant producers of iodine today are Chile and Japan. Due to its
high atomic number and ease of attachment to organic compounds, it has
also found favour as a non-toxic radiocontrast material. Because of
the specificity of its uptake by the human body, radioactive isotopes
of iodine can also be used to treat thyroid cancer. Iodine is also
used as a catalyst in the industrial production of acetic acid and
some polymers.
It is on the World Health Organization's List of Essential Medicines.
History
======================================================================
In 1811, iodine was discovered by French chemist Bernard Courtois, who
was born to a family of manufacturers of saltpetre (an essential
component of gunpowder). At the time of the Napoleonic Wars, saltpetre
was in great demand in France. Saltpetre produced from French nitre
beds required sodium carbonate, which could be isolated from seaweed
collected on the coasts of Normandy and Brittany. To isolate the
sodium carbonate, seaweed was burned and the ashes washed with water.
While investigating the cause of corrosion to the copper vessels used
in the process, Courtois added an excess of sulfuric acid to the waste
remaining and a cloud of violet vapour arose. He noted that the vapour
crystallised on cold surfaces, forming dark crystals. Courtois
suspected that this material was a new element but lacked funding to
pursue it further.
Courtois gave samples to his friends, Charles Bernard Desormes
(1777-1838) and Nicolas Clément (1779-1841), to continue research. He
also gave some of the substance to chemist Joseph Louis Gay-Lussac
(1778-1850), and to physicist André-Marie Ampère (1775-1836). On 29
November 1813, Desormes and Clément made Courtois' discovery public by
describing the substance to a meeting of the Imperial Institut de
France.Desormes and Clément made their announcement at the Institut
impérial de France on 29 November 1813; a summary of their
announcement appeared in the 'Gazette nationale ou Le Moniteur
Universel' of 2 December 1813. See:
*
* On 6 December 1813, Gay-Lussac found and announced that the new
substance was either an element or a compound of oxygen and he found
that it is an element. Gay-Lussac suggested the name "iode"
(anglicised as "iodine"), from the Ancient Greek (, "violet"),
because of the colour of iodine vapour. Ampère had given some of his
sample to British chemist Humphry Davy (1778-1829), who experimented
on the substance and noted its similarity to chlorine and also found
it as an element. Davy sent a letter dated 10 December to the Royal
Society stating that he had identified a new element called iodine.
Arguments erupted between Davy and Gay-Lussac over who identified
iodine first, but both scientists found that both of them identified
iodine first and also knew that Courtois is the first one to isolate
the element.
In 1873, the French medical researcher Casimir Davaine (1812-1882)
discovered the antiseptic action of iodine. Antonio Grossich
(1849-1926), an Istrian-born surgeon, was among the first to use
sterilisation of the operative field. In 1908, he introduced tincture
of iodine as a way to rapidly sterilise the human skin in the surgical
field.
In early periodic tables, iodine was often given the symbol 'J', for
'Jod', its name in German; in German texts, 'J' is still frequently
used in place of 'I'.
Properties
======================================================================
Iodine is the fourth halogen, being a member of group 17 in the
periodic table, below fluorine, chlorine, and bromine; since astatine
and tennessine are radioactive, iodine is the heaviest stable halogen.
Iodine has an electron configuration of [Kr]5s24d105p5, with the seven
electrons in the fifth and outermost shell being its valence
electrons. Like the other halogens, it is one electron short of a full
octet and is hence an oxidising agent, reacting with many elements in
order to complete its outer shell, although in keeping with periodic
trends, it is the weakest oxidising agent among the stable halogens:
it has the lowest electronegativity among them, just 2.66 on the
Pauling scale (compare fluorine, chlorine, and bromine at 3.98, 3.16,
and 2.96 respectively; astatine continues the trend with an
electronegativity of 2.2). Elemental iodine hence forms diatomic
molecules with chemical formula I2, where two iodine atoms share a
pair of electrons in order to each achieve a stable octet for
themselves; at high temperatures, these diatomic molecules reversibly
dissociate a pair of iodine atoms. Similarly, the iodide anion, I−, is
the strongest reducing agent among the stable halogens, being the most
easily oxidised back to diatomic I2. (Astatine goes further, being
indeed unstable as At− and readily oxidised to At0 or At+.)
The halogens darken in colour as the group is descended: fluorine is a
very pale yellow, chlorine is greenish-yellow, bromine is
reddish-brown, and iodine is violet.
Elemental iodine is slightly soluble in water, with one gram
dissolving in 3450 mL at 20 °C and 1280 mL at 50 °C; potassium iodide
may be added to increase solubility via formation of triiodide ions,
among other polyiodides. Nonpolar solvents such as hexane and carbon
tetrachloride provide a higher solubility. Polar solutions, such as
aqueous solutions, are brown, reflecting the role of these solvents as
Lewis bases; on the other hand, nonpolar solutions are violet, the
color of iodine vapour. Charge-transfer complexes form when iodine is
dissolved in polar solvents, hence changing the colour. Iodine is
violet when dissolved in carbon tetrachloride and saturated
hydrocarbons but deep brown in alcohols and amines, solvents that form
charge-transfer adducts.
The melting and boiling points of iodine are the highest among the
halogens, conforming to the increasing trend down the group, since
iodine has the largest electron cloud among them that is the most
easily polarised, resulting in its molecules having the strongest Van
der Waals interactions among the halogens. Similarly, iodine is the
least volatile of the halogens, though the solid still can be observed
to give off purple vapour. Due to this property iodine is commonly
used to demonstrate sublimation directly from solid to gas, which
gives rise to a misconception that it does not melt in atmospheric
pressure. Because it has the largest atomic radius among the halogens,
iodine has the lowest first ionisation energy, lowest electron
affinity, lowest electronegativity and lowest reactivity of the
halogens.
The interhalogen bond in diiodine is the weakest of all the halogens.
As such, 1% of a sample of gaseous iodine at atmospheric pressure is
dissociated into iodine atoms at 575 °C. Temperatures greater than 750
°C are required for fluorine, chlorine, and bromine to dissociate to a
similar extent. Most bonds to iodine are weaker than the analogous
bonds to the lighter halogens. Gaseous iodine is composed of I2
molecules with an I-I bond length of 266.6 pm. The I-I bond is one of
the longest single bonds known. It is even longer (271.5 pm) in solid
orthorhombic crystalline iodine, which has the same crystal structure
as chlorine and bromine. (The record is held by iodine's neighbour
xenon: the Xe-Xe bond length is 308.71 pm.) As such, within the iodine
molecule, significant electronic interactions occur with the two
next-nearest neighbours of each atom, and these interactions give
rise, in bulk iodine, to a shiny appearance and semiconducting
properties. Iodine is a two-dimensional semiconductor with a band gap
of 1.3 eV (125 kJ/mol): it is a semiconductor in the plane of its
crystalline layers and an insulator in the perpendicular direction.
Isotopes
==========
Of the forty known isotopes of iodine, only one occurs in nature,
iodine-127. The others are radioactive and have half-lives too short
to be primordial. As such, iodine is both monoisotopic and
mononuclidic and its atomic weight is known to great precision, as it
is a constant of nature.
The longest-lived of the radioactive isotopes of iodine is iodine-129,
which has a half-life of 15.7 million years, decaying via beta decay
to stable xenon-129. Some iodine-129 was formed along with iodine-127
before the formation of the Solar System, but it has by now completely
decayed away, making it an extinct radionuclide. Its former presence
may be determined from an excess of its daughter xenon-129, but early
attempts to use this characteristic to date the supernova source for
elements in the Solar System are made difficult by alternative nuclear
processes giving iodine-129 and by iodine's volatility at higher
temperatures. Due to its mobility in the environment iodine-129 has
been used to date very old groundwaters. Traces of iodine-129 still
exist today, as it is also a cosmogenic nuclide, formed from cosmic
ray spallation of atmospheric xenon: these traces make up 10−14 to
10−10 of all terrestrial iodine. It also occurs from open-air nuclear
testing, and is not hazardous because of its very long half-life, the
longest of all fission products. At the peak of thermonuclear testing
in the 1960s and 1970s, iodine-129 still made up only about 10−7 of
all terrestrial iodine.
[
http://www.scopenvironment.org/downloadpubs/scope50 SCOPE 50 -
Radioecology after Chernobyl] , the Scientific Committee on Problems
of the Environment (SCOPE), 1993. See table 1.9 in Section 1.4.5.2.
Excited states of iodine-127 and iodine-129 are often used in
Mössbauer spectroscopy.
The other iodine radioisotopes have much shorter half-lives, no longer
than days. Some of them have medical applications involving the
thyroid gland, where the iodine that enters the body is stored and
concentrated. Iodine-123 has a half-life of thirteen hours and decays
by electron capture to tellurium-123, emitting gamma radiation; it is
used in nuclear medicine imaging, including single photon emission
computed tomography (SPECT) and X-ray computed tomography (X-Ray CT)
scans. Iodine-125 has a half-life of fifty-nine days, decaying by
electron capture to tellurium-125 and emitting low-energy gamma
radiation; the second-longest-lived iodine radioisotope, it has uses
in biological assays, nuclear medicine imaging and in radiation
therapy as brachytherapy to treat a number of conditions, including
prostate cancer, uveal melanomas, and brain tumours. Finally,
iodine-131, with a half-life of eight days, beta decays to an excited
state of stable xenon-131 that then converts to the ground state by
emitting gamma radiation. It is a common fission product and thus is
present in high levels in radioactive fallout. It may then be absorbed
through contaminated food, and will also accumulate in the thyroid. As
it decays, it may cause damage to the thyroid. The primary risk from
exposure to high levels of iodine-131 is the chance occurrence of
radiogenic thyroid cancer in later life. Other risks include the
possibility of non-cancerous growths and thyroiditis.
Protection usually used against the negative effects of iodine-131 is
by saturating the thyroid gland with stable iodine-127 in the form of
potassium iodide tablets, taken daily for optimal prophylaxis.
However, iodine-131 may also be used for medicinal purposes in
radiation therapy for this very reason, when tissue destruction is
desired after iodine uptake by the tissue. Iodine-131 is also used as
a radioactive tracer.
Chemistry and compounds
======================================================================
Halogen bond energies (kJ/mol)
X XX HX BX3 AlX3 CX4
F 159 574 645 582 456
Cl |243 |428 |444 |427 |327
Br |193 |363 |368 |360 |272
I |151 |294 |272 |285 |239
Iodine is quite reactive, but it is less so than the lighter halogens,
and it is a weaker oxidant. For example, it does not halogenate carbon
monoxide, nitric oxide, and sulfur dioxide, which chlorine does. Many
metals react with iodine. For the same reason, however, since iodine
has the lowest ionisation energy among the halogens and is the most
easily oxidised of them, it has a more significant cationic chemistry
and its higher oxidation states are rather more stable than those of
bromine and chlorine, for example in iodine heptafluoride.
Charge-transfer complexes
===========================
The iodine molecule, I2, dissolves in CCl4 and aliphatic hydrocarbons
to give bright violet solutions. In these solvents the absorption band
maximum occurs in the 520 – 540 nm region and is assigned to a * to
'σ'* transition. When I2 reacts with Lewis bases in these solvents a
blue shift in I2 peak is seen and the new peak (230 – 330 nm) arises
that is due to the formation of adducts, which are referred to as
charge-transfer complexes.
Hydrogen iodide
=================
The simplest compound of iodine is hydrogen iodide, HI. It is a
colourless gas that reacts with oxygen to give water and iodine.
Although it is useful in iodination reactions in the laboratory, it
does not have large-scale industrial uses, unlike the other hydrogen
halides. Commercially, it is usually made by reacting iodine with
hydrogen sulfide or hydrazine:
:2 I2 + N2H4 4 HI + N2
At room temperature, it is a colourless gas, like all of the hydrogen
halides except hydrogen fluoride, since hydrogen cannot form strong
hydrogen bonds to the large and only mildly electronegative iodine
atom. It melts at −51.0 °C and boils at −35.1 °C. It is an endothermic
compound that can exothermically dissociate at room temperature,
although the process is very slow unless a catalyst is present: the
reaction between hydrogen and iodine at room temperature to give
hydrogen iodide does not proceed to completion. The H-I bond
dissociation energy is likewise the smallest of the hydrogen halides,
at 295 kJ/mol.
Aqueous hydrogen iodide is known as hydroiodic acid, which is a strong
acid. Hydrogen iodide is exceptionally soluble in water: one litre of
water will dissolve 425 litres of hydrogen iodide, and the saturated
solution has only four water molecules per molecule of hydrogen
iodide. Commercial so-called "concentrated" hydroiodic acid usually
contains 48-57% HI by mass; the solution forms an azeotrope with
boiling point 126.7 °C at 56.7 g HI per 100 g solution. Hence
hydroiodic acid cannot be concentrated past this point by evaporation
of water. Unlike gaseous hydrogen iodide, hydroiodic acid has major
industrial use in the manufacture of acetic acid by the Cativa
process.
Other binary iodine compounds
===============================
With the exception of the noble gases, nearly all elements on the
periodic table up to einsteinium (EsI3 is known) are known to form
binary compounds with iodine. Until 1990, nitrogen triiodide was only
known as an ammonia adduct. Ammonia-free NI3 was found to be isolable
at -196 °C but spontaneously decomposes at 0 °C. For thermodynamic
reasons related to electronegativity of the elements, neutral sulfur
and selenium iodides that are stable at room temperature are also
nonexistent, although S2I2 and SI2 are stable up to 183 and 9 K,
respectively. As of 2022, no neutral binary selenium iodide has been
unambiguously identified (at any temperature). Sulfur-iodine and
selenium-iodine polyatomic cations (e.g., [S2I42+][AsF6-]2 and
[Se2I42+][Sb2F11-]2) have been prepared and characterised
crystallographically.
Given the large size of the iodide anion and iodine's weak oxidising
power, high oxidation states are difficult to achieve in binary
iodides, the maximum known being in the pentaiodides of niobium,
tantalum, and protactinium. Iodides can be made by reaction of an
element or its oxide, hydroxide, or carbonate with hydroiodic acid,
and then dehydrated by mildly high temperatures combined with either
low pressure or anhydrous hydrogen iodide gas. These methods work best
when the iodide product is stable to hydrolysis. Other syntheses
include high-temperature oxidative iodination of the element with
iodine or hydrogen iodide, high-temperature iodination of a metal
oxide or other halide by iodine, a volatile metal halide, carbon
tetraiodide, or an organic iodide. For example, molybdenum(IV) oxide
reacts with aluminium(III) iodide at 230 °C to give molybdenum(II)
iodide. An example involving halogen exchange is given below,
involving the reaction of tantalum(V) chloride with excess
aluminium(III) iodide at 400 °C to give tantalum(V) iodide:
3TaCl5 + \underset{(excess)}{5AlI3} -> 3TaI5 + 5AlCl3
Lower iodides may be produced either through thermal decomposition or
disproportionation, or by reducing the higher iodide with hydrogen or
a metal, for example:
TaI5{} + Ta ->[\text{thermal gradient}] [\ce{630^\circ C\ ->\
575^\circ C}] Ta6I14
Most metal iodides with the metal in low oxidation states (+1 to +3)
are ionic. Nonmetals tend to form covalent molecular iodides, as do
metals in high oxidation states from +3 and above. Both ionic and
covalent iodides are known for metals in oxidation state +3 (e.g.
scandium iodide is mostly ionic, but aluminium iodide is not). Ionic
iodides MI'n' tend to have the lowest melting and boiling points among
the halides MX'n' of the same element, because the electrostatic
forces of attraction between the cations and anions are weakest for
the large iodide anion. In contrast, covalent iodides tend to instead
have the highest melting and boiling points among the halides of the
same element, since iodine is the most polarisable of the halogens
and, having the most electrons among them, can contribute the most to
van der Waals forces. Naturally, exceptions abound in intermediate
iodides where one trend gives way to the other. Similarly,
solubilities in water of predominantly ionic iodides (e.g. potassium
and calcium) are the greatest among ionic halides of that element,
while those of covalent iodides (e.g. silver) are the lowest of that
element. In particular, silver iodide is very insoluble in water and
its formation is often used as a qualitative test for iodine.
Iodine halides
================
The halogens form many binary, diamagnetic interhalogen compounds with
stoichiometries XY, XY3, XY5, and XY7 (where X is heavier than Y), and
iodine is no exception. Iodine forms all three possible diatomic
interhalogens, a trifluoride and trichloride, as well as a
pentafluoride and, exceptionally among the halogens, a heptafluoride.
Numerous cationic and anionic derivatives are also characterised, such
as the wine-red or bright orange compounds of and the dark brown or
purplish black compounds of I2Cl+. Apart from these, some
pseudohalides are also known, such as cyanogen iodide (ICN), iodine
thiocyanate (ISCN), and iodine azide (IN3).
Iodine monofluoride (IF) is unstable at room temperature and
disproportionates very readily and irreversibly to iodine and iodine
pentafluoride, and thus cannot be obtained pure. It can be synthesised
from the reaction of iodine with fluorine gas in
trichlorofluoromethane at −45 °C, with iodine trifluoride in
trichlorofluoromethane at −78 °C, or with silver(I) fluoride at 0 °C.
Iodine monochloride (ICl) and iodine monobromide (IBr), on the other
hand, are moderately stable. The former, a volatile red-brown
compound, was discovered independently by Joseph Louis Gay-Lussac and
Humphry Davy in 1813-1814 not long after the discoveries of chlorine
and iodine, and it mimics the intermediate halogen bromine so well
that Justus von Liebig was misled into mistaking bromine (which he had
found) for iodine monochloride. Iodine monochloride and iodine
monobromide may be prepared simply by reacting iodine with chlorine or
bromine at room temperature and purified by fractional
crystallisation. Both are quite reactive and attack even platinum and
gold, though not boron, carbon, cadmium, lead, zirconium, niobium,
molybdenum, and tungsten. Their reaction with organic compounds
depends on conditions. Iodine chloride vapour tends to chlorinate
phenol and salicylic acid, since when iodine chloride undergoes
homolytic fission, chlorine and iodine are produced and the former is
more reactive. However, iodine chloride in carbon tetrachloride
solution results in iodination being the main reaction, since now
heterolytic fission of the I-Cl bond occurs and I+ attacks phenol as
an electrophile. However, iodine monobromide tends to brominate phenol
even in carbon tetrachloride solution because it tends to dissociate
into its elements in solution, and bromine is more reactive than
iodine. When liquid, iodine monochloride and iodine monobromide
dissociate into and ions (X = Cl, Br); thus they are significant
conductors of electricity and can be used as ionising solvents.
Iodine trifluoride (IF3) is an unstable yellow solid that decomposes
above −28 °C. It is thus little-known. It is difficult to produce
because fluorine gas would tend to oxidise iodine all the way to the
pentafluoride; reaction at low temperature with xenon difluoride is
necessary. Iodine trichloride, which exists in the solid state as the
planar dimer I2Cl6, is a bright yellow solid, synthesised by reacting
iodine with liquid chlorine at −80 °C; caution is necessary during
purification because it easily dissociates to iodine monochloride and
chlorine and hence can act as a strong chlorinating agent. Liquid
iodine trichloride conducts electricity, possibly indicating
dissociation to and ions.
Iodine pentafluoride (IF5), a colourless, volatile liquid, is the most
thermodynamically stable iodine fluoride, and can be made by reacting
iodine with fluorine gas at room temperature. It is a fluorinating
agent, but is mild enough to store in glass apparatus. Again, slight
electrical conductivity is present in the liquid state because of
dissociation to and . The pentagonal bipyramidal iodine heptafluoride
(IF7) is an extremely powerful fluorinating agent, behind only
chlorine trifluoride, chlorine pentafluoride, and bromine
pentafluoride among the interhalogens: it reacts with almost all the
elements even at low temperatures, fluorinates Pyrex glass to form
iodine(VII) oxyfluoride (IOF5), and sets carbon monoxide on fire.
Iodine oxides and oxoacids
============================
Iodine oxides are the most stable of all the halogen oxides, because
of the strong I-O bonds resulting from the large electronegativity
difference between iodine and oxygen, and they have been known for the
longest time. The stable, white, hygroscopic iodine pentoxide (I2O5)
has been known since its formation in 1813 by Gay-Lussac and Davy. It
is most easily made by the dehydration of iodic acid (HIO3), of which
it is the anhydride. It will quickly oxidise carbon monoxide
completely to carbon dioxide at room temperature, and is thus a useful
reagent in determining carbon monoxide concentration. It also oxidises
nitrogen oxide, ethylene, and hydrogen sulfide. It reacts with sulfur
trioxide and peroxydisulfuryl difluoride (S2O6F2) to form salts of the
iodyl cation, [IO2]+, and is reduced by concentrated sulfuric acid to
iodosyl salts involving [IO]+. It may be fluorinated by fluorine,
bromine trifluoride, sulfur tetrafluoride, or chloryl fluoride,
resulting iodine pentafluoride, which also reacts with iodine
pentoxide, giving iodine(V) oxyfluoride, IOF3. A few other less stable
oxides are known, notably I4O9 and I2O4; their structures have not
been determined, but reasonable guesses are IIII(IVO3)3 and
[IO]+[IO3]− respectively.
Standard reduction potentials for aqueous I species
!! (acid)!!!! (base)
|I2/I− +0.535 |I2/I− +0.535
|HOI/I− +0.987 IO−/I− +0.48
|0 0 /I− +0.26
|HOI/I2 +1.439 IO−/I2 +0.42
|/I2 +1.195 0 0
|/HOI +1.134 /IO− +0.15
|/ +1.653 0 0
|H5IO6/ +1.601 / +0.65
More important are the four oxoacids: hypoiodous acid (HIO), iodous
acid (HIO2), iodic acid (HIO3), and periodic acid (HIO4 or H5IO6).
When iodine dissolves in aqueous solution, the following reactions
occur:
{{block indent|}}
Hypoiodous acid is unstable to disproportionation. The hypoiodite ions
thus formed disproportionate immediately to give iodide and iodate:
Iodous acid and iodite are even less stable and exist only as a
fleeting intermediate in the oxidation of iodide to iodate, if at all.
Iodates are by far the most important of these compounds, which can be
made by oxidising alkali metal iodides with oxygen at 600 °C and high
pressure, or by oxidising iodine with chlorates. Unlike chlorates,
which disproportionate very slowly to form chloride and perchlorate,
iodates are stable to disproportionation in both acidic and alkaline
solutions. From these, salts of most metals can be obtained. Iodic
acid is most easily made by oxidation of an aqueous iodine suspension
by electrolysis or fuming nitric acid. Iodate has the weakest
oxidising power of the halates, but reacts the quickest.
Many periodates are known, including not only the expected tetrahedral
, but also square-pyramidal , octahedral orthoperiodate ,
[IO3(OH)3]2−, [I2O8(OH2)]4−, and . They are usually made by oxidising
alkaline sodium iodate electrochemically (with lead(IV) oxide as the
anode) or by chlorine gas:
They are thermodymically and kinetically powerful oxidising agents,
quickly oxidising Mn2+ to permanganate, and cleaving glycols,
α-diketones, α-ketols, α-aminoalcohols, and α-diamines. Orthoperiodate
especially stabilises high oxidation states among metals because of
its very high negative charge of −5. Orthoperiodic acid, H5IO6, is
stable, and dehydrates at 100 °C in a vacuum to Metaperiodic acid,
HIO4. Attempting to go further does not result in the nonexistent
iodine heptoxide (I2O7), but rather iodine pentoxide and oxygen.
Periodic acid may be protonated by sulfuric acid to give the cation,
isoelectronic to Te(OH)6 and , and giving salts with bisulfate and
sulfate.
Polyiodine compounds
======================
When iodine dissolves in strong acids, such as fuming sulfuric acid, a
bright blue paramagnetic solution including cations is formed. A
solid salt of the diiodine cation may be obtained by oxidising iodine
with antimony pentafluoride:
{{block indent|2 I2 + 5 SbF5 2 I2Sb2F11 + SbF3}}
The salt I2Sb2F11 is dark blue, and the blue tantalum analogue
I2Ta2F11 is also known. Whereas the I-I bond length in I2 is 267 pm,
that in is only 256 pm as the missing electron in the latter has been
removed from an antibonding orbital, making the bond stronger and
hence shorter. In fluorosulfuric acid solution, deep-blue reversibly
dimerises below −60 °C, forming red rectangular diamagnetic . Other
polyiodine cations are not as well-characterised, including bent
dark-brown or black and centrosymmetric 'C'2'h' green or black ,
known in the and salts among others.
The only important polyiodide anion in aqueous solution is linear
triiodide, . Its formation explains why the solubility of iodine in
water may be increased by the addition of potassium iodide solution:
Many other polyiodides may be found when solutions containing iodine
and iodide crystallise, such as , , , and , whose salts with large,
weakly polarising cations such as Cs+ may be isolated.
Organoiodine compounds
========================
Organoiodine compounds have been fundamental in the development of
organic synthesis, such as in the Hofmann elimination of amines, the
Williamson ether synthesis, the Wurtz coupling reaction, and in
Grignard reagents.
The carbon-iodine bond is a common functional group that forms part of
core organic chemistry; formally, these compounds may be thought of as
organic derivatives of the iodide anion. The simplest organoiodine
compounds, alkyl iodides, may be synthesised by the reaction of
alcohols with phosphorus triiodide; these may then be used in
nucleophilic substitution reactions, or for preparing Grignard
reagents. The C-I bond is the weakest of all the carbon-halogen bonds
due to the minuscule difference in electronegativity between carbon
(2.55) and iodine (2.66). As such, iodide is the best leaving group
among the halogens, to such an extent that many organoiodine compounds
turn yellow when stored over time due to decomposition into elemental
iodine; as such, they are commonly used in organic synthesis, because
of the easy formation and cleavage of the C-I bond. They are also
significantly denser than the other organohalogen compounds thanks to
the high atomic weight of iodine. A few organic oxidising agents like
the iodanes contain iodine in a higher oxidation state than −1, such
as 2-iodoxybenzoic acid, a common reagent for the oxidation of
alcohols to aldehydes, and iodobenzene dichloride (PhICl2), used for
the selective chlorination of alkenes and alkynes. One of the more
well-known uses of organoiodine compounds is the so-called iodoform
test, where iodoform (CHI3) is produced by the exhaustive iodination
of a methyl ketone (or another compound capable of being oxidised to a
methyl ketone), as follows:
Some drawbacks of using organoiodine compounds as compared to
organochlorine or organobromine compounds is the greater expense and
toxicity of the iodine derivatives, since iodine is expensive and
organoiodine compounds are stronger alkylating agents. For example,
iodoacetamide and iodoacetic acid denature proteins by irreversibly
alkylating cysteine residues and preventing the reformation of
disulfide linkages.
Halogen exchange to produce iodoalkanes by the Finkelstein reaction is
slightly complicated by the fact that iodide is a better leaving group
than chloride or bromide. The difference is nevertheless small enough
that the reaction can be driven to completion by exploiting the
differential solubility of halide salts, or by using a large excess of
the halide salt. In the classic Finkelstein reaction, an alkyl
chloride or an alkyl bromide is converted to an alkyl iodide by
treatment with a solution of sodium iodide in acetone. Sodium iodide
is soluble in acetone and sodium chloride and sodium bromide are not.
The reaction is driven toward products by mass action due to the
precipitation of the insoluble salt.
Occurrence and production
======================================================================
Iodine is the least abundant of the stable halogens, comprising only
0.46 parts per million of Earth's crustal rocks (compare: fluorine:
544 ppm, chlorine: 126 ppm, bromine: 2.5 ppm) making it the 60th most
abundant element. Iodide minerals are rare, and most deposits that are
concentrated enough for economical extraction are iodate minerals
instead. Examples include lautarite, Ca(IO3)2, and dietzeite,
7Ca(IO3)2·8CaCrO4. These are the minerals that occur as trace
impurities in the caliche, found in Chile, whose main product is
sodium nitrate. In total, they can contain at least 0.02% and at most
1% iodine by mass. Sodium iodate is extracted from the caliche and
reduced to iodide by sodium bisulfite. This solution is then reacted
with freshly extracted iodate, resulting in comproportionation to
iodine, which may be filtered off.
The caliche was the main source of iodine in the 19th century and
continues to be important today, replacing kelp (which is no longer an
economically viable source), but in the late 20th century brines
emerged as a comparable source. The Japanese Minami Kantō gas field
east of Tokyo and the American Anadarko Basin gas field in northwest
Oklahoma are the two largest such sources. The brine is hotter than 60
°C from the depth of the source. The brine is first purified and
acidified using sulfuric acid, then the iodide present is oxidised to
iodine with chlorine. An iodine solution is produced, but is dilute
and must be concentrated. Air is blown into the solution to evaporate
the iodine, which is passed into an absorbing tower, where sulfur
dioxide reduces the iodine. The hydrogen iodide (HI) is reacted with
chlorine to precipitate the iodine. After filtering and purification
the iodine is packed.
These sources ensure that Chile and Japan are the largest producers of
iodine today. Alternatively, the brine may be treated with silver
nitrate to precipitate out iodine as silver iodide, which is then
decomposed by reaction with iron to form metallic silver and a
solution of iron(II) iodide. The iodine is then liberated by
displacement with chlorine.
Applications
======================================================================
About half of all produced iodine goes into various organoiodine
compounds, another 15% remains as the pure element, another 15% is
used to form potassium iodide, and another 15% for other inorganic
iodine compounds. Among the major uses of iodine compounds are
catalysts, animal feed supplements, stabilisers, dyes, colourants and
pigments, pharmaceutical, sanitation (from tincture of iodine), and
photography; minor uses include smog inhibition, cloud seeding, and
various uses in analytical chemistry.
X-ray imaging
===============
As an element with high electron density and atomic number, iodine
efficiently absorbs X-rays. X-ray radiocontrast agents is the top
application for iodine. In this application, Organoiodine compounds
are injected intravenously. This application is often in conjunction
with advanced X-ray techniques such as angiography and CT scanning. At
present, all water-soluble radiocontrast agents rely on
iodine-containing compounds.
Iodine absorbs X-rays with energies less than 33.3 keV due to the
photoelectric effect of the innermost electrons.
Biocide
=========
Use of iodine as a biocide represents a major application of the
element, ranked 2nd by weight. Elemental iodine (I2) is used as an
antiseptic in medicine. A number of water-soluble compounds, from
triiodide (I3−, generated 'in situ' by adding iodide to poorly
water-soluble elemental iodine) to various iodophors, slowly decompose
to release I2 when applied.
Optical polarising films
==========================
Thin-film-transistor liquid crystal displays rely on polarisation.
The liquid crystal transistor is sandwiched between two polarising
films and illuminated from behind. The two films prevent light
transmission unless the transistor in the middle of the sandwich
rotates the light. Iodine-impregnated polymer films are used in
polarising optical components with the highest transmission and degree
of polarisation.
Co-catalyst
=============
Another significant use of iodine is as a cocatalyst for the
production of acetic acid by the Monsanto and Cativa processes. In
these technologies, hydroiodic acid converts the methanol feedstock
into methyl iodide, which undergoes carbonylation. Hydrolysis of the
resulting acetyl iodide regenerates hydroiodic acid and gives acetic
acid. The majority of acetic acid is produced by these approaches.
Nutrition
===========
Salts of iodide and iodate are used extensively in human and animal
nutrition. This application reflects the status of iodide as an
essential element, being required for two hormones. The production of
ethylenediamine dihydroiodide, provided as a nutritional supplement
for livestock, consumes a large portion of available iodine. Iodine is
a component of iodised salt.
A saturated solution of potassium iodide is used to treat acute
thyrotoxicosis. It is also used to block uptake of iodine-131 in the
thyroid gland (see isotopes section above), when this isotope is used
as part of radiopharmaceuticals (such as iobenguane) that are not
targeted to the thyroid or thyroid-type tissues.
Others
========
Inorganic iodides find specialised uses. Titanium, zirconium, hafnium,
and thorium are purified by the Van Arkel-de Boer process, which
involves the reversible formation of the tetraiodides of these
elements. Silver iodide is a major ingredient to traditional
photographic film. Thousands of kilograms of silver iodide are used
annually for cloud seeding to induce rain.
The organoiodine compound erythrosine is an important food colouring
agent. Perfluoroalkyl iodides are precursors to important surfactants,
such as perfluorooctanesulfonic acid.
(125)I is used as the radiolabel in investigating which ligands go to
which plant pattern recognition receptors (PRRs).
An iodine based thermochemical cycle has been evaluated for hydrogen
production using energy from nuclear power. The cycle has three
steps. At 120 °C, iodine reacts with sulfur dioxide and water to give
hydrogen iodide and sulfuric acid:
:
After a separation stage, at 830 - sulfuric acid splits in sulfur
dioxide and oxygen:
:
Hydrogen iodide, at 300 -, gives hydrogen and the initial element,
iodine:
:
The yield of the cycle (ratio between lower heating value of the
produced hydrogen and the consumed energy for its production, is
approximately 38%. , the cycle is not a competitive means of producing
hydrogen.
Spectroscopy
======================================================================
The spectrum of the iodine molecule, I2, consists of (not exclusively)
tens of thousands of sharp spectral lines in the wavelength range
500-700 nm. It is therefore a commonly used wavelength reference
(secondary standard). By measuring with a spectroscopic Doppler-free
technique while focusing on one of these lines, the hyperfine
structure of the iodine molecule reveals itself. A line is now
resolved such that either 15 components (from even rotational quantum
numbers, 'J'even), or 21 components (from odd rotational quantum
numbers, 'J'odd) are measurable.
Caesium iodide and thallium-doped sodium iodide are used in crystal
scintillators for the detection of gamma rays. The efficiency is high
and energy dispersive spectroscopy is possible, but the resolution is
rather poor.
Chemical analysis
======================================================================
The iodide and iodate anions can be used for quantitative volumetric
analysis, for example in iodometry. Iodine and starch form a blue
complex, and this reaction is often used to test for either starch or
iodine and as an indicator in iodometry. The iodine test for starch is
still used to detect counterfeit banknotes printed on
starch-containing paper.
The iodine value is the mass of iodine in grams that is consumed by
100 grams of a chemical substance typically fats or oils. Iodine
numbers are often used to determine the amount of unsaturation in
fatty acids. This unsaturation is in the form of double bonds, which
react with iodine compounds.
Potassium tetraiodomercurate(II), K2HgI4, is also known as Nessler's
reagent. It is once was used as a sensitive spot test for ammonia.
Similarly, Mayer's reagent (potassium tetraiodomercurate(II) solution)
is used as a precipitating reagent to test for alkaloids. Aqueous
alkaline iodine solution is used in the iodoform test for methyl
ketones.
Biological role
======================================================================
Iodine is an essential element for life and, at atomic number 'Z' =
53, is the heaviest element commonly needed by living organisms.
(Lanthanum and the other lanthanides, as well as tungsten with 'Z' =
74 and uranium with 'Z' = 92, are used by a few microorganisms.) It is
required for the synthesis of the growth-regulating thyroid hormones
tetraiodothyronine and triiodothyronine (T4 and T3 respectively, named
after their number of iodine atoms). A deficiency of iodine leads to
decreased production of T3 and T4 and a concomitant enlargement of the
thyroid tissue in an attempt to obtain more iodine, causing the
disease goitre. The major form of thyroid hormone in the blood is
tetraiodothyronine (T4), which has a longer life than triiodothyronine
(T3). In humans, the ratio of T4 to T3 released into the blood is
between 14:1 and 20:1. T4 is converted to the active T3 (three to four
times more potent than T4) within cells by deiodinases (5'-iodinase).
These are further processed by decarboxylation and deiodination to
produce iodothyronamine (T1a) and thyronamine (T0a'). All three
isoforms of the deiodinases are selenium-containing enzymes; thus
metallic selenium is needed for triiodothyronine and
tetraiodothyronine production.
Iodine accounts for 65% of the molecular weight of T4 and 59% of T3.
Fifteen to 20 mg of iodine is concentrated in thyroid tissue and
hormones, but 70% of all iodine in the body is found in other tissues,
including mammary glands, eyes, gastric mucosa, thymus, cerebrospinal
fluid, choroid plexus, arteries, cervix, salivary glands. During
pregnancy, the placenta is able to store and accumulate iodine. In the
cells of those tissues, iodine enters directly by sodium-iodide
symporter (NIS). The action of iodine in mammal tissues is related to
fetal and neonatal development, and in the other tissues, it is known.
Dietary recommendations and intake
====================================
The daily levels of intake recommended by the United States National
Academy of Medicine are between 110 and 130 μg for infants up to 12
months, 90 μg for children up to eight years, 130 μg for children up
to 13 years, 150 μg for adults, 220 μg for pregnant women and 290 μg
for lactating women. The Tolerable Upper Intake Level (TUIL) for
adults is 1,100 μg/day. This upper limit was assessed by analysing the
effect of supplementation on thyroid-stimulating hormone.
The European Food Safety Authority (EFSA) refers to the collective set
of information as Dietary Reference Values, with Population Reference
Intake (PRI) instead of RDA, and Average Requirement instead of EAR;
AI and UL are defined the same as in the United States. For women and
men ages 18 and older, the PRI for iodine is set at 150 μg/day; the
PRI during pregnancy and lactation is 200 μg/day. For children aged
1-17 years, the PRI increases with age from 90 to 130 μg/day. These
PRIs are comparable to the U.S. RDAs with the exception of that for
lactation.
The thyroid gland needs 70 μg/day of iodine to synthesise the
requisite daily amounts of T4 and T3. The higher recommended daily
allowance levels of iodine seem necessary for optimal function of a
number of body systems, including mammary glands, gastric mucosa,
salivary glands, brain cells, choroid plexus, thymus, arteries.
Natural food sources of iodine include seafood which contains fish,
seaweeds, kelp, shellfish and other foods which contain dairy
products, eggs, meats, vegetables, so long as the animals ate iodine
richly, and the plants are grown on iodine-rich soil. Iodised salt is
fortified with potassium iodate, a salt of iodine, potassium, oxygen.
As of 2000, the median intake of iodine from food in the United States
was 240 to 300 μg/day for men and 190 to 210 μg/day for women. The
general US population has adequate iodine nutrition, with lactating
women and pregnant women having a mild risk of deficiency. In Japan,
consumption was considered much higher, ranging between 5,280 μg/day
to 13,800 μg/day from wakame and kombu that are eaten, both in the
form of kombu and wakame and kombu and wakame umami extracts for soup
stock and potato chips. However, new studies suggest that Japan's
consumption is closer to 1,000-3,000 μg/day. The adult UL in Japan was
last revised to 3,000 μg/day in 2015.
After iodine fortification programs such as iodisation of salt have
been done, some cases of iodine-induced hyperthyroidism have been
observed (so-called Jod-Basedow phenomenon). The condition occurs
mainly in people above 40 years of age, and the risk is higher when
iodine deficiency is high and the first rise in iodine consumption is
high.
Deficiency
============
In areas where there is little iodine in the diet, which are remote
inland areas and faraway mountainous areas where no iodine rich foods
are eaten, iodine deficiency gives rise to hypothyroidism, symptoms of
which are extreme fatigue, goitre, mental slowing, depression, low
weight gain, and low basal body temperatures. Iodine deficiency is the
leading cause of preventable intellectual disability, a result that
occurs primarily when babies or small children are rendered
hypothyroidic by no iodine. The addition of iodine to salt has largely
destroyed this problem in wealthier areas, but iodine deficiency
remains a serious public health problem in poorer areas today. Iodine
deficiency is also a problem in certain areas of all continents of the
world. Information processing, fine motor skills, and visual problem
solving are normalised by iodine repletion in iodine-deficient people.
Toxicity
==========
{{Chembox
| container_only = yes
| Section1 =
}}
Elemental iodine (I2) is toxic if taken orally undiluted. The lethal
dose for an adult human is 30 mg/kg, which is about 2.1-2.4 grams for
a human weighing 70 to 80 kg (even when experiments on rats
demonstrated that these animals could survive after eating a 14000
mg/kg dose and are still living after that). Excess iodine is more
cytotoxic in the presence of selenium deficiency. Iodine
supplementation in selenium-deficient populations is problematic for
this reason. The toxicity derives from its oxidising properties,
through which it denaturates proteins (including enzymes).
Elemental iodine is also a skin irritant. Solutions with high
elemental iodine concentration, such as tincture of iodine and Lugol's
solution, are capable of causing tissue damage if used in prolonged
cleaning or antisepsis; similarly, liquid Povidone-iodine (Betadine)
trapped against the skin resulted in chemical burns in some reported
cases.
Occupational exposure
=======================
The U.S. Occupational Safety and Health Administration (OSHA) has set
the legal limit (Permissible exposure limit) for iodine exposure in
the workplace at 0.1 ppm (1 mg/m3) during an 8-hour workday. The
National Institute for Occupational Safety and Health (NIOSH) has set
a Recommended exposure limit (REL) of 0.1 ppm (1 mg/m3) during an
8-hour workday. At levels of 2 ppm, iodine is immediately dangerous to
life and health.
Allergic reactions
====================
Some people develop a hypersensitivity to products and foods
containing iodine. Applications of tincture of iodine or Betadine can
cause rashes, sometimes severe. Parenteral use of iodine-based
contrast agents (see above) can cause reactions ranging from a mild
rash to fatal anaphylaxis. Such reactions have led to the
misconception (widely held, even among physicians) that some people
are allergic to iodine itself; even allergies to iodine-rich foods
have been so construed. In fact, there has never been a confirmed
report of a true iodine allergy, as an allergy to iodine or iodine
salts is biologically impossible. Hypersensitivity reactions to
products and foods containing iodine are apparently related to their
other molecular components; thus, a person who has demonstrated an
allergy to one food or product containing iodine may not have an
allergic reaction to another. Patients with various food allergies
(fishes, shellfish, eggs, milk, seaweeds, kelp, meats, vegetables,
kombu, wakame) do not have an increased risk for a contrast medium
hypersensitivity. The patient's allergy history is relevant.
US DEA List I status
======================
Phosphorus reduces iodine to hydroiodic acid, which is a reagent
effective for reducing ephedrine and pseudoephedrine to
methamphetamine. For this reason, iodine was designated by the United
States Drug Enforcement Administration as a List I precursor chemical
under 21 CFR 1310.02.
License
=========
All content on Gopherpedia comes from Wikipedia, and is licensed under CC-BY-SA
License URL:
http://creativecommons.org/licenses/by-sa/3.0/
Original Article:
http://en.wikipedia.org/wiki/Iodine
